1. Production
    In the laboratory  From alum schists  From clays or bauxite  From cryolite 

2. Uses

3. Chemical reactions

4. References
    Footnotes  Notations 

5. External links

Aluminium sulfate is a chemical compound with the formula Al₂(SO₄)₃. It is soluble in water and is mainly used as a coagulating agent (promoting particle collision by neutralizing charge) in the purification of drinking water^[3][4] and waste water treatment plants, and also in paper manufacturing.

The anhydrous form occurs naturally as a rare mineral millosevichite, found e.g. in volcanic environments and on burning coal-mining waste dumps. Aluminium sulfate is rarely, if ever, encountered as the anhydrous salt. It forms a number of different hydrates, of which the hexadecahydrate Al₂(SO₄)₃•16H₂O and octadecahydrate Al₂(SO₄)₃•18H₂O are the most common. The heptadecahydrate, whose formula can be written as [Al(H₂O)₆]₂(SO₄)₃•5H₂O, occurs naturally as the mineral alunogen.

Aluminium sulfate is sometimes called alum or papermaker's alum in certain industries. However, the name "alum" is more commonly and properly used for any double sulfate salt with the generic formula {{chem|XAl(SO|4|)|2|·12H|2|O}}, where X is a monovalent cation such as potassium or ammonium.^[5]

Production

In the laboratory

Aluminium sulfate may be made by adding aluminium hydroxide, Al(OH)₃, to sulfuric acid, H₂SO₄:

2 Al(OH)₃ + 3 H₂SO₄ → Al₂(SO₄)₃ + 6H₂O

or by heating aluminum metal in a sulfuric acid solution:

2 Al + 3 H₂SO₄ → Al₂(SO₄)₃ + 3 H₂↑

From alum schists

The alum schists employed in the manufacture of aluminium sulfate are mixtures of iron pyrite, aluminium silicate and various bituminous substances, and are found in upper Bavaria, Bohemia, Belgium, and Scotland. These are either roasted or exposed to the weathering action of the air. In the roasting process, sulfuric acid is formed and acts on the clay to form aluminium sulfate, a similar condition of affairs being produced during weathering. The mass is now systematically extracted with water, and a solution of aluminium sulfate of specific gravity 1.16 is prepared. This solution is allowed to stand for some time (in order that any calcium sulfate and basic ferric sulfate may separate), and is then evaporated until ferrous sulfate crystallizes on cooling; it is then drawn off and evaporated until it attains a specific gravity of 1.40. It is now allowed to stand for some time, and decanted from any sediment.{{sfn|Chisholm|1911|p=767}}

From clays or bauxite

In the preparation of aluminum sulfate from clays or from bauxite, the material is gently calcined, then mixed with sulfuric acid and heated gradually to boiling; it is allowed to stand for some time, the clear solution drawn off.

From cryolite

When cryolite is used as the ore, it is mixed with calcium carbonate and heated. By this means, sodium aluminate is formed; it is then extracted with water and precipitated either by sodium bicarbonate or by passing a current of carbon dioxide through the solution. The precipitate is then dissolved in sulfuric acid.{{sfn|Chisholm|1911|p=767}}

Uses

It is sometimes used in the human food industry as a firming agent for food starch, where it takes on E number E520, and in animal feed as a bactericide. Aluminum sulfate may be used as a deodorant, an astringent, or as a stiptic for superficial shaving wounds.^[6]

It improves vaccine immunogenicity as a vaccine adjuvant "by facilitating the slow release of antigen from the vaccine depot formed at the site of inoculation."^[6]

Aluminium sulfate is used in water purification and as a mordant in dyeing and printing textiles. In water purification, it causes suspended impurities to coagulate into larger particles and then settle to the bottom of the container (or be filtered out) more easily. This process is called coagulation or flocculation. Research suggests that in Australia, aluminium sulfate used this way in drinking water treatment is the primary source of hydrogen sulfide gas in sanitary sewer systems.^[7] An improper and excess application incident in 1988 polluted the water supply of Camelford in Cornwall.

When dissolved in a large amount of neutral or slightly alkaline water, aluminium sulfate produces a gelatinous precipitate of aluminium hydroxide, Al(OH)₃. In dyeing and printing cloth, the gelatinous precipitate helps the dye adhere to the clothing fibers by rendering the pigment insoluble.

Aluminium sulfate is sometimes used to reduce the pH of garden soil, as it hydrolyzes to form the aluminium hydroxide precipitate and a dilute sulfuric acid solution. An example of what changing the pH level of soil can do to plants is visible when looking at Hydrangea macrophylla. The gardener can add aluminium sulfate to the soil to reduce the pH which in turn will result in the flowers of the Hydrangea turning a different color (blue). The aluminium is what makes the flowers blue; at a higher pH, the aluminium is not available to the plant.^[8]

In the construction industry, it is used as waterproofing agent and accelerator in concrete. Another use is a foaming agent in fire fighting foam.

In limnology (lake & pond science), Alum is often effectively employed to control levels of soluble ortho-phosphate (PO4) and to agglomerate suspended dirt particles for improved clarification. The product can be applied as a granular product or as "Liquid Alum" which is water + dissolved ALum (48.6% Alum dry basis). pH must be monitored in lakes with valuable fish when considering an Alum treatment. Alum is quite acidic and a lake with low pH (below 7) or a lake lacking a natural carbonate buffering capacity may find that the pH falls below 6 and to level when fish as harmed or killed. In such cases, a buffering agent - such as lime or Soda Ash or Sodium Bicarbonate - may be applied simultaneously with the Alum to prevent a precipitous drop in pH.

It can also be very effective as a molluscicide,^[9] killing spanish slugs.

{{chem|Al|2|(SO|4|)|3}}   +   3 {{chem|Pb(CH|3|CO|2|)|2}}   →   2 {{chem|Al(CH|3|CO|2|)|3}}   +   3 {{chem|PbSO|4}}

{{chem|Al|2|(SO|4|)|3}}   +   2 {{chem|Pb(CH|3|CO|2|)|2}}   →   {{chem|Al|2|SO|4|(CH|3|CO|2|)|4}}   +   2 {{chem|PbSO|4}}

Chemical reactions

The compound decomposes to γ−alumina and sulfur trioxide when heated between 580 and 900 °C. It combines with water forming hydrated salts of various compositions.

Aluminium sulfate reacts with sodium bicarbonate to which foam stabilizer has been added, producing carbon dioxide for fire-extinguishing foams:

Al₂(SO₄)₃ + 6 NaHCO₃ → 3 Na₂SO₄ + 2 Al(OH)₃ + 6 CO₂

The carbon dioxide is trapped by the foam stabilizer and creates a thick foam which will float on top of hydrocarbon fuels and seal off access to atmospheric oxygen, smothering the fire. Chemical foam was unsuitable for use on polar solvents such as alcohol, as the fuel would mix with and break down the foam blanket. The carbon dioxide generated also served to propel the foam out of the container, be it a portable fire extinguisher or fixed installation using hoselines. Chemical foam is considered obsolete in the United States and has been replaced by synthetic mechanical foams, such as AFFF which have a longer shelf life, are more effective, and more versatile, although some countries such as Japan and India continue to use it.{{citation needed|date=August 2010}}

References

Footnotes

Notations

• {{cite book | author=Pauling, Linus | title=General Chemistry | location=W.H. Freeman | publisher=San Francisco | year=1970 | isbn=978-0-486-65622-9}}

External links

• International Chemical Safety Card 1191
• [https://www.cdc.gov/niosh/npg/npgd0024.html NIOSH Pocket Guide to Chemical Hazards]
• WHO Food Additive Series No. 12
• Aluminum and health
• Government of Canada Fact Sheets and Frequently Asked Questions: Aluminum Salts

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